1. The balanced chemical equation provides information about the beginning and end of reaction. The reaction mechanism gives the path of the reaction. Mechanisms provide a very detailed picture of which bonds are broken and formed during the course of a reaction. Elementary Steps Elementary step: any process that occurs in a single step. Reaction Mechanisms

2. Elementary Steps Molecularity: the number of molecules present in an elementary step.  Unimolecular: one molecule in the elementary step,  Bimolecular: two molecules in the elementary step, and  Termolecular: three molecules in the elementary step. It is not common to see termolecular processes (statistically improbable). Reaction Mechanisms

3. Multistep Mechanisms Some reaction proceed through more than one step: NO 2 (g) + NO 2 (g)  NO 3 (g) + NO(g) NO 3 (g) + CO(g)  NO 2 (g) + CO 2 (g) Notice that if we add the above steps, we get the overall reaction: NO 2 (g) + CO(g)  NO(g) + CO 2 (g) Reaction Mechanisms

4. Multistep Mechanisms If a reaction proceeds via several elementary steps, then the elementary steps must add to give the balanced chemical equation. Intermediate: a species which appears in an elementary step which is not a reactant or product. Reaction Mechanisms

5. Rate Laws for Elementary Steps The rate law of an elementary step is determined by its molecularity: –Unimolecular processes are first order, –Bimolecular processes are second order, and –Termolecular processes are third order. Rate Laws for Multistep Mechanisms Rate-determining step: is the slowest of the elementary steps. Reaction Mechanisms

6. Rate Laws for Elementary Steps Reaction Mechanisms

7. Rate Laws for Multistep Mechanisms Therefore, the rate-determining step governs the overall rate law for the reaction. Mechanisms with an Initial Fast Step It is possible for an intermediate to be a reactant. Consider 2NO(g) + Br 2 (g)  2NOBr(g) Reaction Mechanisms

8. Mechanisms with an Initial Fast Step 2NO(g) + Br 2 (g)  2NOBr(g) The experimentally determined rate law is Rate = k[NO] 2 [Br 2 ] Consider the following mechanism Reaction Mechanisms

9. Mechanisms with an Initial Fast Step The rate law is (based on Step 2): Rate = k 2 [NOBr 2 ][NO] The rate law should not depend on the concentration of an intermediate (intermediates are usually unstable). Assume NOBr 2 is unstable, so we express the concentration of NOBr 2 in terms of NOBr and Br 2 assuming there is an equilibrium in step 1 we have Reaction Mechanisms

10. Mechanisms with an Initial Fast Step By definition of equilibrium: Therefore, the overall rate law becomes Note the final rate law is consistent with the experimentally observed rate law. Reaction Mechanisms

11. A catalyst changes the rate of a chemical reaction. There are two types of catalyst: –homogeneous, and –heterogeneous. Chlorine atoms are catalysts for the destruction of ozone. Homogeneous Catalysis The catalyst and reaction is in one phase. Catalysis

12. Catalysis

13. Homogeneous Catalysis Hydrogen peroxide decomposes very slowly: 2H 2 O 2 (aq)  2H 2 O(l) + O 2 (g) In the presence of the bromide ion, the decomposition occurs rapidly: –2Br - (aq) + H 2 O 2 (aq) + 2H + (aq)  Br 2 (aq) + 2H 2 O(l). –Br 2 (aq) is brown. –Br 2 (aq) + H 2 O 2 (aq)  2Br - (aq) + 2H + (aq) + O 2 (g). –Br - is a catalyst because it can be recovered at the end of the reaction. Catalysis

14. Homogeneous Catalysis Generally, catalysts operate by lowering the activation energy for a reaction. Catalysis